A sample of hydrogen (H2) gas was collected over water at 60°C. If the total volume of gas collected was 55.6 mL and the atmospheric pressure was 1.1 bar, what mass (in g) of hydrogen gas was collected above the water? The vapour pressure of water at 60°C is 150 torr.

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KendyI423210 KendyI423210 · Answer 1 · 2022-10-26T14:13:38+02:00

First, calculate the partial pressure of H2 by considering the total pressure:

[tex]P_{\text{tot}}=P_{H2}+P_{H20}[/tex]

where,

Ptot = 1.1bar

PH20 = 150Torr

Convert the previous values to atm:

Ptot = 1.1 bar = 1.085 atm

PH20 = 150 Torr = 0.197 atm

Then, for PH2 you obtain:

[tex]P_{H2}=P_{\text{tot}}-P_{H20}=1.085\text{atm-}0.197\text{atm}=0.888atm[/tex]

Next, use the following equation for ideal gases:

[tex]PV=\text{nRT}[/tex]

where,

V: volume = 55.65mL = 0.05565 L

R: gas constant = 0.082 atm*L/mol*K

T: temperature = 60 + 273 = 333K

P: partial pressure of H2 = 0.888atm

Solve the equation for n and replace the values of the other parameters:

[tex]n=\frac{PV}{RT}=\frac{(0.888atm)(0.05565L)}{(0.082atm\cdot\frac{L}{\text{mol}\cdot K})(333K)}=0.00181mol[/tex]

Next, use the atomic weight to determine the mass of H2:

[tex]M=(\frac{2.0141g}{\text{mol}})(0.00181mol)=0.0036g[/tex]

Hence, there are 0.0036 g of H2 collected above the water